Presentation on theme: "Atomic Radius By: Kelly Sun, David Zeng, George Xu."— Presentation transcript:
Atomic Radius By: Kelly Sun, David Zeng, George Xu
Definition of Atomic Radius Atomic Radius is a measure of the size of an atom Total distance from atoms nucleus to the outermost electron orbital Usually found be finding half the distance between two nuclei. No standard definition There are four types of radii: covalent atomic radii, ionic radii, metallic radii, and van der Waals radii Measured in picometers or angstroms
Covalent Atomic Radii For atoms that are bonded by covalent bonds and share electrons, the radius of the atom is less than if the atoms were just touching each other or tangent to each other. Covalent atomic radii are measured for nonmetallic atoms
Ionic Radius The radius of an atom forming an ionic bond or ion. Element’s share of the distance between neighboring ions in an ionic solid. Generally: Cations are smaller than their parent ions Anions are larger than their parent ions
Metallic Radii Metallic Radii are the radii of atoms joined by a metallic bond (similar to covalent bonds)
van der Waals Radii van der Waals radii apply for the most part to only noble gases The two atoms are just touching because the attractive forces bonding them are much less, leaving the two atoms “unsquashed”
Description of Trend The atomic radius decreases across a period: moving from left to right along the periodic table The atomic radius increases down a group Atomic Radius is not a fixed property as electrons do not move in a fixed orbit within their orbitals
Reasons for Trends The atomic radius increases down a group: Why: As the number of energy levels increase moving down a group, so does the number of electrons present Each subsequent energy level is further from the nucleus than the last. Therefore, the atomic radius increases as the group and energy levels increase. The shielding effect also increases the distance electrons are to the nucleus as inner electrons have a repulsive charge on the outer electrons. The atomic radius decreases across a period: Why: Moving across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge” meaning there is a stronger force pulling the electrons closer to the nucleus resulting in a small atomic radius.
How it’s measured (Determination) Large degree of uncertainty since they do not have definite orbits No fixed radius Covalent Radii are found by measuring bond lengths between pairs of covalently bonded atoms and dividing the length by two Metallic Radii are found in a way similar to covalent radii but between pairs of atoms bonded by metallic bonds Van-der-Waals radii are determined from the contact distances between unbonded atoms in touching molecules or atoms. Ionic radii are found by using X-Ray crystallography- Crystalline atoms cause a beam of x-rays to diffract into many specific directions; the angles are measured, and a three dimensional picture of structure and density can be found.
Anomalies Some Noble Gases, although the farthest atoms to the right in the group 8A column, are larger than atoms in the group 7A column. This is because atomic radius can only be measured by taking the distance between bonded atoms and Noble Gases do not form covalent and ionic bonds due to their full electron shell. As a result, atomic radius for these can be only estimated using Van Der Waals forces which is not as strong. Example: Neon, Radon In the Lanthanide Series, the atomic radius decreases at a greater rate from left to right than expected. Normally, as nucleus becomes more charged across a period, the decrease in atomic radius is partially counteracted by a shielding effect. However, in the lanthanide series, the f shell electrons is not effective at shielding the outer shells. As a result, atomic radius decreases more quickly across the series.
Anomalies 2 Transition Metals do not exhibit the general trend as the rest of the periodic table. Moving across the periods on the periodic table for transition metals shows an increase in atomic radius rather than a decrease. This due to to the ability of transitional metals to accept unpaired electrons from other atoms to fill vacancies in their d orbitals, which would therefore increase the atomic radius.
Example Problems Order elements from smallest to largest: Ne, F, Ba, Na, Mg, H, Al